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Redox Reactions of Transition Metals

Applied Science: BTEC Level 3 National Extended Diploma, topic 1.2.13

Oxidation of Transition Metals

Transition metal ions can undergo oxidation reactions. Remember that an "oxidation is a loss" of electrons.

Oxidation of Fe2+ to Fe3+

Oxidation of Fe2+ to Fe3+

  • Iron (II) ions can be oxidised by potassium manganate (VII) solution in acidic conditions.
  • The observed colour change is from green to brown.
  • Iron is oxidised from oxidation state +2 to +3.
  • Manganese is reduced from oxidation state +7 to +2.
Half-equation to equation

Half-equation to equation

  • You need to understand how to construct an equation from the half-equations. Consider these two half-equations:
  • Fe2+ → Fe3+ + e-
    MnO4- + 8H+ + 5e- → Mn2+ + 4H2O
  • One e- is produced with Fe3+ and five e- are in the reactants of the MnO4- half-equation.
  • You therefore need five times more of the first half-equation.
  • MnO4- + 8H+ + 5Fe2+ → Mn2+ + 4H2O + 5Fe3+
Oxidation of Cr3+ to Cr2O72-

Oxidation of Cr3+ to Cr2O72-

  • Cr3+ in an excess of OH- ions (from NaOH) exists in the complex ion [Cr(OH)6]3-.
  • This complex ion and be oxidised with H2O2:
  • 2[ Cr(OH)6 ]3- + 3H2O2 → 2CrO42- + 2OH- + 8H2O
  • The colour changes observed are from purple to green when excess NaOH is added and from green to yellow when H2O2 is added.
CrO42- to Cr2O72-

CrO42- to Cr2O72-

  • A reversible reaction can be used to easily change between CrO42- and Cr2O72-:
  • CrO42-  ⇌add  OH−add  H+  \;\xrightleftharpoons[add\; OH^-]{add\;H^+}\;Cr2O72-
  • Use dilute sulphuric acid and sodium hydroxide to push the equlibrium to either side.
Colour change and oxidation states

Colour change and oxidation states

  • The colour change obeserved is from yellow (CrO42-) to orange (Cr2O72-).
  • Chromium is oxidised from oxidation state +3 to +6.
  • Oxygen is reduced from oxidation state -1 to -2.

Reduction of Transition Metals

Transition metal ions can undergo reduction reactions. Remember that a "reduction is a gain" of electrons.

Reduction of iron (III) to iron (II)

Reduction of iron (III) to iron (II)

  • Iron (III) ions can be reduced by using aqueous iodide ions as the reducing agent.
  • The observed colour change is from brown to green.
The equation

The equation

  • Fe3+ + 2I- → 2Fe2+ + I2
  • Iron is reduced from oxidation state +3 to +2.
  • Iodine is oxidised from oxidation state -1 to 0.
Reduction of chromium (VI) to chromium (III)

Reduction of chromium (VI) to chromium (III)

  • Cr2O72- can be reducted with zinc in an acidic solution.
  • Cr2O72- + 14H+ + 3Zn → 2Cr3+ + 7H2O + 3Zn2+
  • A colour change from orange to green is observed.
  • Chromium is reduced from oxidation state +6 to +3.
  • Zinc is oxidised from oxidation state 0 to +2.
Reduction of copper (II) to copper (I)

Reduction of copper (II) to copper (I)

  • Cu2+ can be reduced by iodide ions. In the process, the iodide ions are oxidised by the copper (II).
  • 2Cu2+(aq) + 4I- → 2CuI(s) + I2 (aq)
Colour change and oxidation states

Colour change and oxidation states

  • A colour change from orange to brown. An off-white precipitate of copper (I) iodide settles in the solution over time.
  • Copper is reduced from oxidation state +2 to +1.
  • Iodine is oxidised from oxidation state -1 to 0.

Disproportionation

Disproportionation occurs when an element both oxidises and reduces in a reaction.

Disproportionation of copper (I)

Disproportionation of copper (I)

  • In solution, copper (I) ions (oxidation state +1) forms solid copper (oxidation state 0) and copper (II) ions (oxidation state +2).
  • It is therefore both reduced and oxidised in the solution.
  • 2Cu+(aq) → Cu(s) + Cu2+(aq)
Example of disproportionation

Example of disproportionation

  • Copper (I) oxide can be dissolved in dilute sulphuric acid in a disproportionation reaction:
  • Cu2O(s) + H2SO4 (s) → Cu(s) + CuSO4 (aq) H2O(l)
  • Copper goes from oxidation state +1 in Cu2O(s) to oxidation state 0 in Cu(s) and oxidation state +2 in CuSO4 (aq).
  • The copper (I) has been both reduced and oxidised.

Test yourself on Redox Reactions of Transition Metals

  • Iron (II) is oxidised to iron (III) in a reaction with potassium manganate (VII) solution in acidic conditions. What colour change is observed in this reaction?

  • Chromium (III) is oxidised to chromium (VI) in a reaction with excess sodium hydroxide and hydrogen peroxide. What colour change is observed in this reaction?

  • What is the observed colour change when iron (III) is reduced to iron (II)?

  • What is the observed colour change when chromium (VI) is reduced to chromium (III)?

  • What is a disproportionation reaction?

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